Ionic Bonding
AdvancedMetal atoms transfer outer electrons to non-metal atoms, forming cations and anions; the electrostatic attraction between them is the ionic bond. Ionic compounds have giant ionic structures — high melting points, no conduction as solids, conduction when molten or dissolved.
- The motive for bonding: a stable outermost shell — the octet, or helium's 2
- Metals lose electrons to form cations; non-metals gain to form anions; electron transfer creates the ionic bond (electrostatic attraction)
- Write formulas by charge balance: NaCl, MgCl₂, CaO, Al₂O₃
文體小知識：科學知識說明文——以電子轉移解釋離子鍵的形成，並連繫離子化合物的巨型結構與高熔點、導電性等性質。
§Text Passage
1Every atom's chemical goal is stability: a full outermost shell of 8 electrons — the octet (helium manages with 2) — matching a noble gas. Metal atoms have few outer electrons, only 1 to 3, and lose them easily; non-metal atoms have many, 5 to 7, and readily gain electrons to fill up. One loses, the other gains — a perfect match.
2Take the formation of sodium chloride: a sodium atom (2,8,1) transfers its outermost electron to a chlorine atom (2,8,7), becoming a sodium ion Na⁺ (2,8) with charge +1; the chlorine atom becomes a chloride ion Cl⁻ (2,8,8) with charge −1. Both ions now have the same stable arrangements as the noble gases neon and argon, and the two oppositely charged particles attract each other.
3The electrostatic attraction between cation and anion is the ionic bond. Sodium chloride consists of countless Na⁺ and Cl⁻ alternating in a giant ionic structure: each ion is surrounded by six ions of the opposite charge, with attraction acting in every direction. Breaking this structure takes huge energy, so ionic compounds have high melting points — sodium chloride melts at 801°C. The solids are hard but brittle, shattering along planes when struck.
4Ionic conductivity depends on state: in a solid the ions are locked in fixed lattice positions and cannot move, so it does not conduct. When molten or dissolved in water, the ions swim freely and the substance conducts. Formulas such as NaCl, MgCl₂ and CaO reflect the ratios in which the charges cancel — the rule for writing the formulas of ionic compounds.
1Summary 重點摘要
Metal atoms transfer outer electrons to non-metal atoms, forming cations and anions; the electrostatic attraction between them is the ionic bond. Ionic compounds have giant ionic structures — high melting points, no conduction as solids, conduction when molten or dissolved.
2Vocabulary cards Vocabulary
The electrostatic attraction between cations and anions that holds an ionic compound together.
An ion formed by losing electrons and carrying positive charge, e.g. Na⁺.
An ion formed by gaining electrons and carrying negative charge, e.g. Cl⁻.
The handing-over of electrons from metal atoms to non-metal atoms — the key step in forming ionic bonds.
The attraction between opposite charges — the nature of the ionic bond.
The huge regular lattice of alternating cations and anions, as in sodium chloride.
The temperature at which a solid melts; ionic compounds have high values, e.g. NaCl at 801°C.
The ability to carry current; ionic compounds conduct only when molten or dissolved.
A diagram using dots and crosses to show outermost electrons and how bonding occurs.
The regular, repeating three-dimensional arrangement of ions.
3Step by step Steps
- Step 1: The octet is the motiveThe great motive for bonding is reaching a noble-gas-like stable outermost shell: 8 electrons (helium manages with 2) — the octet. For example, sodium at 2,8,1 needs only to lose 1 electron to become 2,8, matching neon; chlorine at 2,8,7 needs only 1 more electron to become 2,8,8, matching argon. Who tends to lose and who tends to gain is decided by the outermost electron count — the first step in understanding ionic bonding and a direct application of last lesson's periodic table work.
- Step 2: Metals lose electrons to form cationsMetal atoms hold only 1 to 3 outer electrons; once these are gone, the inner shells are already stable. So metals tend to lose electrons, become positively charged, and form cations. For example, sodium (2,8,1) loses 1 electron to become Na⁺ (2,8); magnesium (2,8,2) loses 2 to become Mg²⁺; calcium (2,8,8,2) loses 2 to become Ca²⁺. In the ion's symbol the magnitude goes before the sign — Ca²⁺, never reversed — and the positive charge equals the number of electrons lost.
- Step 3: Non-metals gain electrons to form anionsNon-metal atoms carry 5 to 7 outer electrons, only one to three short of a full shell, so they tend to take in electrons, become negatively charged, and form anions. For example, chlorine (2,8,7) gains 1 electron to become Cl⁻ (2,8,8); oxygen (2,6) gains 2 to become O²⁻ (2,8). Both types of ion now match noble-gas arrangements and are stable. In electron diagrams the incoming electrons are drawn with crosses, showing clearly who transferred electrons to whom — the presentation exams expect.
- Step 4: Electron transfer and forming the ionic bondThe formation of sodium chloride is the textbook demonstration: a sodium atom (2,8,1) transfers its outermost electron to a chlorine atom (2,8,7), becoming Na⁺ while the chlorine becomes Cl⁻ — both reach the octet. The oppositely charged Na⁺ and Cl⁻ attract electrostatically, and this attraction is the ionic bond. For example, in an electron diagram draw sodium's electrons as dots and chlorine's as crosses; after the transfer that cross joins chlorine's outermost shell, and the direction of transfer is plain to see.
- Step 5: Formulas and charge balanceIonic formulas follow the rule of charge balance: the positive and negative charges must cancel to zero. For example, Na⁺ with Cl⁻ is 1 to 1, giving NaCl; Mg²⁺ needs 2 Cl⁻ to cancel, giving MgCl₂; Ca²⁺ with O²⁻ is 1 to 1, giving CaO; Al³⁺ needs 3 O²⁻, crossing over to give Al₂O₃. In fact the outermost electron count already hints at the charge: Group I loses 1 to give +1, Group VII gains 1 to give −1. The periodic table backs up every formula — check the group before writing.
- Step 6: Giant ionic structure and high melting pointsIn a sodium chloride crystal, Na⁺ and Cl⁻ alternate in an orderly array, each ion surrounded by six oppositely charged ions — a giant ionic structure, or lattice. Ionic bonds run throughout the crystal, and melting means breaking enormous numbers of strong electrostatic attractions, which takes immense energy. For example, sodium chloride melts at 801°C, so table salt resists an ordinary kitchen flame; ice, held by weak intermolecular forces, melts at 0°C — clear evidence of structural difference. In answers, cover all three points: strong forces, huge numbers, high energy.
- Step 7: The conductivity puzzle of statesSolid ionic compounds do not conduct because the ions are fixed at lattice positions, vibrating but not travelling. When molten or dissolved in water, the ions are free to move toward the electrodes, and the substance conducts. For example, connect solid salt to a battery and bulb and nothing lights up; connect salt solution and the bulb glows — thanks to freely moving Na⁺ and Cl⁻. The answer mantra: ions present and free to move — then it conducts. The contrast among solid, molten and aqueous states is a high-frequency explanation question.
4Diagrams Diagrams
5Worked example Example
Electron transfer and properties: sodium burns in chlorine to form sodium chloride. Explain the formation of Na⁺ and Cl⁻ by electron transfer (giving each electron arrangement), why sodium chloride has a high melting point, and in which states it conducts electricity.
- Electron transfer: sodium (2,8,1) loses its 1 outermost electron to become Na⁺ (2,8); chlorine (2,8,7) gains 1 electron to become Cl⁻ (2,8,8)
- Bonding: oppositely charged Na⁺ and Cl⁻ attract electrostatically — the ionic bond — and alternate in a giant ionic structure
- High melting point: melting requires breaking a huge number of strong electrostatic attractions throughout the lattice, so it takes immense energy — NaCl melts at 801°C
- Conduction: in the solid the ions are fixed and cannot move, so it does not conduct; when molten or dissolved, the ions move freely and it conducts
- Summary: formation is by electron transfer; properties follow from structure — the heart of the bonding-structure-property link
- Answer: A sodium atom (2,8,1) transfers its outermost electron to a chlorine atom (2,8,7). The sodium becomes a Na⁺ ion with charge +1 and arrangement 2,8; the chlorine becomes a Cl⁻ ion with charge −1 and arrangement 2,8,8 — both matching noble-gas structures. The oppositely charged ions attract electrostatically, forming ionic bonds, and alternate in a giant ionic structure. Because these attractions are strong and run throughout the lattice, melting takes immense energy, so sodium chloride melts at 801°C. In the solid the ions are locked in place and it does not conduct; when molten or dissolved in water the ions move freely, so it conducts.
6Common mistakes Mistakes
7Quick check Quick check
Q1What is the nature of the ionic bond?
Q2How do sodium and chlorine atoms become ions?
Q3Why does sodium chloride have a high melting point?
Q4In which states does sodium chloride conduct, and why?
Q5What is the formula of the compound of magnesium (2,8,2) and chlorine?
8Recap Recap
- The motive for bonding: a stable outermost shell — the octet, or helium's 2
- Metals lose electrons to form cations; non-metals gain to form anions; electron transfer creates the ionic bond (electrostatic attraction)
- Write formulas by charge balance: NaCl, MgCl₂, CaO, Al₂O₃
- The giant ionic structure gives high melting points (NaCl 801°C) and hard, brittle solids
- Conduction rule: ions present and free to move — conducts when molten or dissolved, not as a solid